Name: ph at key points in a weak acid titration
Uploaded: 2020-09-24T15:06:51.000Z
Duration: 3 min 55 s
Channel: JoVE

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Titration Calculations: Weak Acid - Strong Base

De pH in verschillende stadia van een zwak zuur of base titratie wordt berekend met verschillende methoden tijdens verschillende punten van de titratie. Als een zwak zuur of zwakke base de belangrijkste bepalende factor is voor de pH, worden de Ka of Kb en een ICE-tabel of de Henderson-Hasselbalch-vergelijking gebruikt. Als er na de neutralisatiereactie sterk zuur of sterke base aanwezig is, wordt de concentratie van de resterende hydronium-of hydroxide-ionen gebruikt om de pH te berekenen.De aanvankelijke pH van 50 milliliter van een 0, 10 molaire azijnzuuroplossing is 2, 87 en wordt berekend met behulp van de Ka en een ICE-tabel, aangezien azijnzuur de belangrijkste bijdrage levert. Als deze oplossing met 0, 0050 mol acetaat wordt getitreerd met 0, 10 molair natriumhydroxide, reageren de hydroxide-ionen met azijnzuur om acetaat te produceren, wat resulteert in een buffer. Daarom wordt, wanneer 10 milliliter natriumhydroxide dat 0, 0010 mol hydroxide-ionen bevat, wordt toegevoegd, 0, 0010 mol acetaat gevormd en 0, 0040 mol azijnzuur blijft achter.De pH van de buffer kan worden berekend door deze waarden in de Henderson-Hasselbalch-vergelijking te vervangen en is gelijk aan 4, 14. Wanneer 25 milliliter natriumhydroxide wordt toegevoegd, wordt de helft van de aanvankelijke molen azijnzuur omgezet in acetaat. Op dit punt is de pH gelijk aan de pKa, aangezien de hoeveelheid azijnzuur en acetaat-ion gelijk zijn.De verdere toevoeging van natriumhydroxide tot 50 milliliter zet alle azijnzuurmoleculen om in acetaat en het equivalentiepunt wordt bereikt. Aangezien acetaationen basisch zijn, ligt het equivalentiepunt in het basische gebied. De concentratie van het acetaation wordt berekend door het aantal mol te delen door het totale volume van de oplossing.De pH wordt bepaald met behulp van de Kb voor acetaationen en een ICE-tabel, aangezien het acetaation de belangrijkste bijdrage levert aan de pH op het equivalentiepunt. De Kb voor acetaat wordt berekend met de formule Kw is gelijk aan Ka maal Kb en is gelijk aan 5, 6 10⁻¹⁰. Het vervangen van evenwichtsconcentraties in de uitdrukking voor Kb geeft de hydroxideconcentratie, 5, 3 x 10⁻⁶ molair.De pOH en pH van de oplossing zijn respectievelijk 5, 28 en 8, 72. Verdere toevoeging van natriumhydroxide aan de oplossing resulteert in een mengsel van acetaationen en natriumhydroxide. De uiteindelijke concentratie natriumhydroxide bepaalt echter de pH van de oplossing, aangezien het een sterkere base is dan acetaat.Daarom, als 70 milliliter natriumhydroxide aan de oplossing wordt toegevoegd, kan de eindconcentratie van de hydroxide-ionen worden berekend door het totale aantal molen azijnzuur, 0, 0050 molen, af te trekken van het totale aantal molen toegevoegde hydroxide-ionen, 0, 0070 molen, en het te delen door het totale volume van de oplossing, 120 milliliter of 0, 12 liter. Aangezien de hydroxide-ionenconcentratie 0, 017 molair is, worden de pOH en pH van de oplossing berekend op respectievelijk 1, 78 en 12, 22.

titratieberekeningen: zwak zuur - sterke base

Berekening van de pH voor titratieoplossingen: zwak zuur/sterke base Voor de titratie van 25,00 ml 0,100 M CH3CO2H met 0,100 M NaOH kan de reactie als volgt worden weergegeven: <img alt="Eq1" src="/files/ftp_upload/11414/11414_Equation_1.png" style="text-align: center; height: 50px;"> De pH van de titratieoplossing na toevoeging van de verschillende volumes NaOH-titrant kan als volgt worden berekend: (a) De initiële pH wordt berekend voor de azijnzuuroplossing volgens de gebruikelijke ICE-benadering: <img alt="Eq2" src="/files/ftp_upload/11414/11414_Equation_2.png" style="text-align: center; height: 155px;"> (b) Het zuur en de titrant zijn beide monoprotisch en de monster- en titrantoplossingen zijn gelijk geconcentreerd; dit volume titrant vertegenwoordigt dus het equivalentiepunt. In tegenstelling tot het sterkzure voorbeeld bevat het reactiemengsel in dit geval een zwakke geconjugeerde base (acetaation). De pH van de oplossing wordt berekend rekening houdend met de base-ionisatie van acetaat, dat aanwezig is in een concentratie van <img alt="Eq3" src="/files/ftp_upload/11414/11414_Equation_3.png" style="text-align: center; height: 72px;"> Base-ionisatie van acetaat wordt weergegeven door de vergelijking <img alt="Eq4" src="/files/ftp_upload/11414/11414_Equation_4.png" style="text-align: center; height: 115px;"> Uitgaande van x << 0,0500, kan de pH worden berekend via de gebruikelijke ICE-benadering: <img alt="Eq5" src="/files/ftp_upload/11414/11414_Equation_5.png" style="text-align: center; height: 150px;"> Merk op dat de pH op het equivalentiepunt van deze titratie aanzienlijk groter is dan 7, zoals verwacht bij het titreren van een zwak zuur met een sterke base. (c) Titrantvolume = 12,50 ml. Dit volume vertegenwoordigt de helft van de stoichiometrische hoeveelheid titrant, en dus is de helft van het azijnzuur geneutraliseerd om een equivalente hoeveelheid acetaation op te leveren. De concentraties van deze geconjugeerde zuur-basepartners zijn daarom gelijk. Een handige benadering voor het berekenen van de pH is het gebruik van de Henderson-Hasselbalch-vergelijking: <img alt="Eq6" src="/files/ftp_upload/11414/11414_Equation_6.png" style="text-align: center; height: 115px;"> (pH = pKa op het half-equivalentiepunt in een titratie van een zwak zuur) (d) Titrantvolume = 37,50 ml. Dit volume vertegenwoordigt een stoichiometrische overmaat titrant en een reactieoplossing die zowel het titratieproduct, het acetaation als de overmaat sterke titrant bevat. In dergelijke oplossingen wordt de pH van de oplossing voornamelijk bepaald door de hoeveelheid overtollige sterke base: <img alt="Eq7" src="/files/ftp_upload/11414/11414_Equation_7.png" style="text-align: center; height: 155px;"> Deze tekst is aangepast van<a href="https://openstax.org/books/chemistry-2e/pages/14-7-acid-base-titrations"> Openstax, Chemistry 2e, Section 14.7: Acid-base Titrations.</a>

Titratieberekeningen: zwak zuur - sterke base

Calculating pH for Titration Solutions: Weak Acid/Strong Base
For the titration of 25.00 mL of 0.100 M CH3CO2H with 0.100 M NaOH, the reaction can be ...

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ph at key points in a weak acid titration

pH at key points in a weak acid titration can be calculated for an acetic acid solution as sodium hydroxide is added. The example uses 25.00 mL of 0.100 M CH3CO2H titrated with 0.100 M NaOH. The reaction is shown as the weak acid reacts with the strong base.

The initial pH is found from the acetic acid solution using the usual ICE approach. ICE means initial, change, and equilibrium, and it is a common way to set up equilibrium problems. The solution pH is calculated before any titrant is added.

When 25.00 mL of NaOH has been added, the titration reaches the equivalence point. The acid and titrant are both monoprotic, and the sample and titrant have the same concentration. At this point, the mixture contains acetate ion, which is the weak conjugate base of acetic acid.

Because acetate ion is present, the pH at the equivalence point is found from its base ionization. The acetate concentration is used in an ICE setup, and x is assumed to be much less than 0.0500. The pH at this point is significantly greater than 7, which is expected for a weak acid titrated with a strong base.

At 12.50 mL of NaOH, the titration is at the half-equivalence point. This is one-half of the stoichiometric amount of titrant, so one-half of the acetic acid has been neutralized. Equal amounts of acetic acid and acetate are present, and the Henderson-Hasselbalch equation is a convenient way to find the pH. At the half-equivalence point, pH equals pKa.

At 37.50 mL of NaOH, there is a stoichiometric excess of titrant. The reaction mixture contains acetate ion from the titration product and extra strong base. In this case, the pH is determined mainly by the amount of excess base.

pH at Key Points in a Weak Acid Titration

ph-at-key-points-in-a-weak-acid-titration

common ion effect

Common Ion Effect on Solubility

The common ion effect reduces the solubility of an ionic compound in a solution that already contains one of its ions. In aqueous solutions, this means less of the solid dissolves than it would in pure water. The effect follows the law of mass action and can be explained with Le Châtelier’s principle. Silver iodide shows this equilibrium clearly. When silver(I) ions or iodide ions are added, the dissolution equilibrium shifts left. More AgI precipitates, and the concentrations of dissolved Ag+...

buffers

Buffer Solutions and pH Stability

Buffer solutions are mixtures that help keep pH steady. A buffer contains an appreciable amount of a weak conjugate acid-base pair. It resists a change in pH when small amounts of a strong acid or a strong base are added. One common buffer uses a weak acid and its salt. Acetic acid and sodium acetate form this type of system, written as CH3COOH(aq) + CH3COONa(aq). Another buffer uses a weak base and its salt. A solution of ammonia and ammonium chloride, NH3(aq) + NH4Cl(aq), is an example. The...

calculating ph changes in a buffer solution

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Buffer pH After Acid or Base Addition

A buffer solution can resist sudden pH changes when a strong acid or strong base is added. It does not stop the pH from changing completely. The shift is usually small, as long as the addition stays within the buffer’s capacity. To find the new pH, first determine how the added acid or base changes the amounts of the buffer components. A buffer contains a weak acid and its conjugate base, or the reverse pair. These amounts are found with stoichiometric calculations, which track how much of...

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Buffer Capacity and pH Stability

Buffer solutions help keep pH fairly constant, but they do not work forever. Their resistance to pH change depends on having enough of a weak acid and its conjugate base, or a weak base and its conjugate acid, in the mixture. When too much strong acid or strong base is added, one part of the buffer pair is used up and the buffer action becomes weaker. Buffer capacity describes how much acid or base can be added to a fixed volume of buffer before the pH changes noticeably, usually by about one...

titration calculations: strong acid - strong base

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pH at Key Points in Acid-Base Titration

pH at key points in a strong acid and strong base titration can be calculated from the amounts of acid and base present. The example uses 25.00 mL of 0.100 M HCl, a strong acid, and 0.100 M NaOH, a strong base. The pH changes as base is added, and each stage follows a different calculation. At 0 mL of titrant, the pH comes from the ionization of HCl. Because HCl is a strong acid, its ionization is complete. The hydronium ion molarity is 0.100 M, and the pH is found from that value. At 12.50...

titration calculations: weak acid - strong base

pH at key points in a weak acid titration can be calculated for an acetic acid solution as sodium hydroxide is added. The example uses 25.00 mL of 0.100 M CH3CO2H titrated with 0.100 M NaOH. The reaction is shown as the weak acid reacts with the strong base. The initial pH is found from the acetic acid solution using the usual ICE approach. ICE means initial, change, and equilibrium, and it is a common way to set up equilibrium problems. The solution pH is calculated before any titrant is...

indicators

Acid-Base Color Changes

Acid-base indicators show color changes that help identify the pH of a solution. These are often weak organic acids or weak organic bases, and they respond to changes in hydronium ion concentration in dilute aqueous solution. Phenolphthalein is one example. It is colorless when the hydronium ion concentration is greater than 5.0 × 10^-9 M, which means the pH is below 8.3. In more basic solutions, where the hydronium ion concentration is lower, it turns red or pink. Methyl orange shows how an...

titration of a polyprotic acid

Polyprotic Acid Equivalence Points

Polyprotic acids have more than one ionizable hydrogen, so they lose protons in steps during titration. When the acid dissociation constants for those protons are far enough apart, the titration curve shows a separate equivalence point for each hydrogen that can ionize. Diprotic acids have two ionizable hydrogens and produce two equivalence points. Triprotic acids have three ionizable hydrogens and produce three equivalence points. Carbonic acid, H2CO3, is a weak diprotic acid. Carbonic acid...

solubility equilibria

Predicting Precipitation with Ksp

Solubility equilibria explain how a dissolved ionic compound and its solid form can exist in balance. In a saturated solution, the solute dissolves and the ions precipitate at equal rates. These ideas matter in real settings such as tooth decay and water purification, where compound solubility affects the outcome. A substance with finite solubility can form a saturated solution when more solute is added than the solvent can dissolve. The result is a heterogeneous mixture of the saturated...

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Common Ion Effects on Solubility

The common ion effect lowers the solubility of an ionic compound in water when the solution already contains one of the ions made by that compound. It follows the law of mass action and can be explained with Le Châtelier’s principle. When a common ion is added, the dissolution equilibrium shifts toward the solid, so less of the salt stays dissolved. Silver iodide shows this effect clearly. If silver ions or iodide ions are added to a solution of AgI, the equilibrium shifts left and more AgI...

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Ligands and Complex Ion Stability

Complex ions form when a central atom, often a transition metal cation, binds to ligands. Ligands are ions or molecules around the metal ion. They can be neutral molecules such as H2O or NH3, or ions such as CN− or OH−. In many cases, the ligands act as Lewis bases by donating a pair of electrons to the central atom. These reactions are part of coordination chemistry, which studies coordination complexes. When a metal ion reacts with one or more ligands to form a complex, the equilibrium...

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Solubility Limits and Selective Precipitation

Solubility limits and selective precipitation explain when ions form a solid in water. The key idea is the balance between an insoluble compound and its dissolved ions. In aqueous solutions that contain calcium ions and carbonate ions, this equilibrium can exist even if the solution was not made by saturating water with calcium carbonate. When two soluble salts are mixed, such as sodium carbonate and calcium nitrate, the ion concentrations determine whether a precipitate forms. If the reaction...

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Cation Identification by Precipitation Tests

Cation identification by precipitation tests uses selective reactions to find metal ions in a mixture. In qualitative analysis, different reagents are added step by step. Each reaction can produce a characteristic precipitate for a specific group of cations. Other steps, such as changing pH, heating the mixture to dissolve a precipitate again, or adding reagents that form complex ions, help separate ions within the same group. Group IV cations such as Ba2+, Ca2+, and Mg2+ are identified...

acid-base titration curves

Tracking pH During Acid-Base Titration

Acid-base titration uses pH changes to track how a solution responds as titrant is added. A titration curve is a graph of a solution property versus the amount of added titrant. In acid-base titrations, pH is the main property to follow because it changes in a predictable way with the solution makeup and helps show the endpoint. This pattern can be seen with a strong acid and strong base, a strong acid and weak base, or a strong base and weak acid. For example, 25.00 mL of 0.100 M HCl titrated...

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Buffer pH with the Henderson-Hasselbalch Equation

The Henderson-Hasselbalch equation helps relate buffer pH to the weak acid in the solution. It comes from the ionization-constant expression for a weak acid. By rearranging that expression to solve for hydronium ion concentration, [H3O+], and then taking the negative logarithm of both sides, the equation can be written in terms of pH. In this form, pKa is defined as the negative logarithm of the weak acid ionization constant, or pKa = −log Ka. The equation shows the connection between pH, pKa,...

key terms and concepts

Follow how pH changes during a weak acid titration with strong base, including the initial solution, half-equivalence point, equivalence point, and excess base.

pH at Key Points in a Weak Acid Titration | JoVE K12

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Temperature Dependence on Reaction Rate

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The Arrhenius equation relates the activation energy and the rate constant, k, for chemical reactions. In the Arrhenius equation, k = Ae&#8722;Ea/RT, R is ...

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Relating Reaction Mechanisms
In a multistep reaction mechanism, one of the elementary steps progresses significantly slower than the others. This slowest ...

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Homogeneous Equilibria for Gaseous Reactions
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Relative Strengths of Conjugate Acid-Base Pairs

Br&#248;nsted-Lowry acid-base chemistry is the transfer of protons; thus, logic suggests a relation between the relative strengths of conjugate acid-base ...

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Strong Acid and Base Solutions

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Le Chatelier's Principle: Changing Concentration

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Weak Base Solutions

Some compounds produce hydroxide ions when dissolved by chemically reacting with water molecules. In all cases, these compounds react only partially and ...

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Le Chatelier's Principle: Changing Volume (Pressure)

For gas-phase equilibria, changes in the concentrations of reactants and products can occur with altered volume and pressure. The partial pressure, P, of ...

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The pH of a solution containing an acid can be determined using its acid dissociation constant and its initial concentration. If a solution contains two ...

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Le Chatelier's Principle: Changing Temperature

Consistent with the law of mass action, an equilibrium stressed by a change in concentration will shift to re-establish equilibrium without any change in ...

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Salts with Acidic Ions
Salts are ionic compounds composed of cations and anions, either of which may be capable of undergoing an acid or base ionization ...

Bronsted-Lowry Acids and Bases

The acid-base reaction class has been studied for quite some time. In 1680, Robert Boyle reported traits of acid solutions that included their ability to ...

Determining the pH of Salt Solutions

The pH of a salt solution is determined by its component anions and cations. Salts that contain pH-neutral anions and the hydronium ion-producing cations ...

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Acid/Base Strengths and Dissociation Constants

The relative strength of an acid or base is the extent to which it ionizes when dissolved in water. If the ionization reaction is essentially complete, ...

Polyprotic Acids

Acids are classified by the number of protons per molecule that they can give up in a reaction. Acids such as HCl, HNO3, and HCN that contain one ...

Chemical reactions
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Content elaboration: In the Physical Science course, coefficients were used to balance simple equations. Other representations, including Lewis structures and three-dimensional models, were also used and manipulated to demonstrate the conservation of matter in chemical reactions. In this course, more complex reactions will be studied, classified and represented with balanced chemical equations and three-dimensional models. Classifying reactions into types can be a helpful organizational tool for recognizing patterns of what may happen when two substances are mixed. Teachers should be aware that the common reaction classifications that are often used in high school chemistry courses may lead to misconceptions because they are not based on the actual chemistry, but on surface features that can be similar from one system to another (e.g., exchanging partners), even though the underlying chemistry is not the same. However, these classifications may be useful in making predictions about what happens when two substances are mixed. Some general types of chemical reactions are oxidation/reduction, synthesis, decomposition, single replacement, double replacement (including precipitation reactions and some acid-base neutralizations) and combustion reactions. Some reactions can fit into more than one category. For example, a single replacement reaction can also be classified as an oxidation/reduction reaction. Identification of reactions involving oxidation and reduction as well as indicating what substance is being oxidized and what is being reduced are appropriate in this course. However, balancing complex oxidation/reduction reactions is reserved for more advanced study. Organic molecules release energy when undergoing combustion reactions and are used to meet the energy needs of society (e.g., oil, gasoline, natural gas) and to provide the energy needs of biological organisms (e.g., cellular respiration). When a reaction between two ionic compounds in aqueous solution results in the formation of a precipitate or molecular compound, the reaction often occurs because the new ionic or covalent bonds are stronger than the original ion-dipole interactions of the ions in solution. Laboratory experiences (3-D or virtual) with different types of chemical reactions should be provided. Reactions occur when reacting particles collide in an appropriate orientation and with sufficient energy. The rate of a chemical reaction is the change in the amount of the reactants or products in a specific period of time. Increasing the probability or effectiveness of the collisions between the particles increases the rate of the reaction. Therefore, changing the concentration of the reactants, changing the temperature or the pressure of gaseous reactants, or using a catalyst, can change the reaction rate. Likewise, the collision theory can be applied to dissolving solids in a liquid solvent and can be used to explain why reactions are more likely to occur between reactants in the aqueous or gaseous state than between solids. The rate at which a substance dissolves should not be confused with the amount of solute that can dissolve in a given amount of solvent (solubility). Mathematical treatment of reaction rates is reserved for more advanced study. Computer simulations can help visualize reactions from the perspective of the kinetic-molecular theory. In middle school, the differences between potential and kinetic energy and the particle nature of thermal energy were introduced. For chemical systems, potential energy is in the form of chemical energy and kinetic energy is in the form of thermal energy. The total amount of chemical energy and/or thermal energy in a system is impossible to measure. However, the energy change of a system can be calculated from measurements (mass and change in temperature) from calorimetry experiments in the laboratory. Conservation of energy is an important component of calorimetry equations. Thermal energy is the energy of a system due to the movement of its particles. The thermal energy of an object depends upon the temperature and chemical composition. Some materials require little energy to change their temperature and other materials require a great deal to change their temperature by the same amount. Specific heat is a measure of how much energy is needed to change the temperature of a specific mass of material a specific amount. Specific heat values can be used to calculate the thermal energy change, the temperature (initial, final or change in) or mass of a material in calorimetry. Water has a particularly high specific heat capacity, which is important in regulating Earth's temperature. As studied in middle school, chemical energy is the potential energy associated with chemical systems. Chemical reactions involve valence electrons forming bonds to yield more stable products with lower energies. Energy is required to break interactions and bonds between the reactant atoms and energy is released when an interaction or bond is formed between the atoms in the products. Molecules with weak bonds (e.g., ATP) are less stable and tend to react to produce more stable products, releasing energy in the process. Generally, energy is transferred out of the system (exothermic) when the products have stronger bonds than the reactants and is transferred into the system (endothermic) when the reactants have stronger bonds than the products. Predictions of the energy requirements (endothermic or exothermic) of a reaction can be made given a table of bond energies. Graphic representations can be drawn and interpreted to represent the energy changes during a reaction. The role of energy in determining the spontaneity of chemical reactions is dealt with conceptually in this course. Entropy and its influence on the spontaneity of reactions are reserved for more advanced study. All reactions are reversible to a degree and many reactions do not proceed completely toward products but appear to stop progressing before the reactants are all used up. At this point, the amounts of the reactants and the products appear to be constant and the reaction can be said to have reached dynamic equilibrium. Dynamic equilibrium means the rate of the reverse reaction is equal to the rate of the forward reaction so there is no apparent change in the reaction. If a chemical system at equilibrium is disturbed by a change in the conditions of the system (e.g., increase or decrease in the temperature, pressure on gaseous equilibrium systems, concentration of a reactant or product), then the equilibrium system will respond by shifting to a new equilibrium state, reducing the effect of the change (Le Chatelier's Principle). If products are removed as they are formed during a reaction, then the equilibrium position of the system is forced to shift to favor the products. In this way, an otherwise unfavorable reaction can be made to occur. Mathematical treatment of equilibrium reactions is reserved for advanced study. Computer simulations can help visualize the progression of a reaction to dynamic equilibrium and the continuation of both the forward and reverse reactions after equilibrium has been attained. Properties of acids and bases and the ranges of the pH scale were introduced in Physical Science. In this course, the structural features of molecules are explored to further understand acids and bases. Acids often result when hydrogen is covalently bonded to an electronegative element and is easily dissociated from the rest of the molecule to bind with water to form a hydronium ion (H3O+). The acidity of an aqueous solution can be expressed as pH, where pH can be calculated from the concentration of the hydronium ion. Bases are likely to dissociate in water to form a hydroxide ion. Acids can react with bases to form a salt and water. Such neutralization reactions can be studied quantitatively by performing titration experiments.

pH at Key Points in a Weak Acid Titration

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